Flame Tests and Atomic Emission Spectroscopy

Objective

In this lab students will learn about atomic energy levels, emission spectroscopy, and flame tests for element identification.

Overview

This lab consists of two parts. In the first part students will draw their own continuous spectrum. Then they will analyze the emission lines from at least four different atomic emission light sources. These light sources are gas discharge tubes filled with gaseous samples of various elements. They will record the spectra they observe in such a way as to relate them to the continuous spectrum they drew.

In the second part students will use small samples of 6 chloride salts of different metals. These they will place into a flame in order to observe the colors produced. These colors come from the excitation of electrons which then resume their ground states by emitting light of very specific colors.

Background

The electrons in an atom occupy different energy levels, as you know. When all of the electrons are at the lowest possible energy level they are said to be in the ground state. Electrons do not always stay in the ground state. Sometimes they can be promoted to a higher-energy electron shell. This can happen in two ways. First, the electron can absorb a photon of just the right amount of energy to move it from one quantum shell to another. Second, when atoms are heated their electrons can gain energy from the heat. This promotes them to the higher-energy shell. When an electron is in a higher-energy shell it is said to be in an excited state.

Electrons in excited states do not usually stay in them for very long. When electrons lose their energy they do so by emitting a photon of light. Photons are particles with energy but no mass. Their energy is directly proportional to the frequency of the light (remember: E = hf). The photons emitted precisely match the quantum energy difference between the excited state and the ground state.

For different elements the spacing between the ground state and the higher energy levels is different. This gives rise to a way to uniquely identify elements based on their spectrum. A spectrum is the scientific name for a rainbow: light broken into the different wavelengths that make it up. You can see spectra using a spectroscope, a prism or a diffraction grating. The back of an ordinary CD is a diffraction grating. Atoms (as opposed to molecules) produce very sharp lines in a spectrum when they are heated. You will look at these lines in Part I of this lab. These lines show the energy differences between the excited states and the ground state. An example of the atomic spectrum of hydrogen is shown below:

When you look at the hydrogen gas discharge tube you will see a mixture of these four colors. To see the lines you have to use a diffraction grating or a prism. Even so, the mixed color alone can be enough to identify an element. Put simply, each element glows a unique color when heated sufficiently. You will use this color to identify elements in this lab when you do flame tests in Part II.

Materials

Safety

• Wear goggles or risk sitting out the lab
• Do not allow chemicals to touch bare skin: wash well with water immediately if you touch anything accidentally
• Use caution with the burner, it is very hot!
• Wash your hands with soap and water after handling the chemicals
• CuCl₂ (Copper(II) Chloride) is highly toxic by ingestion; avoid contact with eyes, skin and mucous membranes.
• HCl (Hydrochloric acid) is extremely corrosive and highly toxic by ingestion; avoid contact with eyes, skin and mucous membranes.
• LiCl (Lithium Chloride) is moderately toxic by ingestion; avoid contact with eyes, skin and mucous membranes.

Procedure

Remember to record your observations in your lab notebook or on a piece of paper in your binder before you leave class. When making observations be sure to use all of your senses except taste. Never taste anything in the chemistry lab. Chances are good you will regret it if you do.

Part I: Observing Atomic Emission Spectra

1. Obtain a clean sheet of plain paper and colored pencils.
2. Near the top of the page (with the long side held vertical) write the title: Atomic Emission Spectra.
3. Use a ruler to draw a rectangle that spans the width of the page, leaving about 1.5 cm for a margin on either side. The rectangle should be about 4 cm high.
4. Make a scale inside the rectangle showing the wavelengths of the visible spectrum. If you wish, you may color your scale according to the table shown at right. Use the scale shown below as a guide.
5. Below your visible spectrum scale draw in four more rectangles of the same size, including the scale.
6. Obtain a CD, a diffraction grating, or a spectroscope. Any of these devices will split the light produced by the elements in the tubes into a spectrum you can see. Observe the spectra of four different elements and draw the spectrum of each one in one of the rectangles on your Atomic Emission Spectra page. Line up the lines as best you as can with the continuous spectrum so that you can estimate the wavelength of the lines you draw.

Part II: Flame Tests

1. Make five little dishes using Al foil. Label the dishes with the name and formula of each of the salts you will be using. This is most easily done by labeling a sheet of scrap paper in a grid with the names and formulas of the compounds meant to go in each dish. Simply keep each dish on its label when not in use.
2. Do not put the copper chloride into an aluminum dish! When wet it reacts with the aluminum and burns a hole in it. Instead, put the copper chloride in a 50 mL beaker.
3. Collect a small sample (less than 0.5 g) of each of the known metal salts which your teacher has provided and carry them all to your lab bench.
4. Obtain an inoculation loop for your group.
5. Obtain 20 - 30 mL of 3.0 M HCl in your 100 mL beaker.
6. Each group member must record information in a neat table with the following columns. Make this table before you even turn on the gas.
   1. Name & Formula of Metal Chloride
   2. Metal Ion
   3. Color of Flame
   4. Approx. Wavelength (nm)
   5. Approx. Wavelength (m)
7. Clean the inoculation loop by swirling it gently in the acid. Be careful, this strong acid can cause severe burns if you get it on your skin. Then, once you light the burner, heat the loop until it glows red hot. This step removes any ions clinging to the loop from previous experiments.
8. Light and adjust your bunsen burner. Be sure to clean your loop carefully.
9. To do a flame test with each metal salt use a pipet to put a few drops of water on your salt sample. Swirl it to dissolve using the inoculation loop. Get a film of the solution inside the loop and bring it into the hottest part of the flame. Repeat the dip into the salt solution as often as necessary to see the flame test color.
10. Carefully note the color of each metal salt when it is put in the flame. Use the chart on the previous page to estimate the approximate wavelength of the color you see. Record all data in the table you made earlier.
11. Clean the inoculation loop using the acid and heating method each time you change from one metal salt to another. Failing to do so will result in mixed flame test colors.
12. Your teacher has prepared two solutions with two of the metal salts. They are labelled Unknown A and Unknown B. Use a 50 mL beaker to get a few mL of Unknown A.
13. Use the inoculation loop to see what flame color the unknown solution produces. Write down the identification of the unknown metal based on your observation.
14. Clean out the beaker using the method recommended by your instructor (hazardous wastes must be disposed of properly). Repeat the above procedure for Unknown B.

Lab Report

For this lab you must turn in the following items:

1. Your Atomic Emission Spectra page from Part I
2. Your data table from Part II: Flame Tests
3. Answers to the following questions
4. A short essay about how fireworks are made, concentrating on how they produce different colors

Questions

1. What is a spectroscope and what is it for? Remember, you used a spectroscope in Part I of this lab.
2. In Part I you observed the spectral lines for four different elements. What is happening within an atom that causes it to emit light in specific lines in a spectrum?
3. Why did each of the different elements have a different emission spectrum? Explain your answer.
4. Could you use the emission spectrum of an element to identify it? How?
5. Carefully determine the wavelength of each of the emission lines in the spectra of two of the elements that you observed. Calculate the frequency (c = λf) and the energy (E = hf) for each of the lines for the two elements. Make a table of your results with the following columns: 1 wavelength in meters, 2 frequency in Hz, and 3 Energy in J. List the results in order from the least energetic to the most energetic photons.
6. Why do different metals have different characteristic flame test colors? (Refer to your results from Part II).
7. Most salts contain a metal and a non-metal. Look at the compounds we tested and explain how we can be sure that it is the metal atoms that are responsible for the colors you see.
8. Could flame tests be useful in determining identities of metals in a mixture? If so, what problems might arise? If not, why not? Explain your answer.
9. What colors did your unknowns produce in the flame? What are your unknowns?
10. Why do the chemicals have to be heated in the flame before the colored light is emitted?
11. In Part I what was it that enabled the gaseous atoms in the discharge tubes to emit light?
12. In your own words, write a short explanation of how an electron absorbs energy and re-emits it as light and why different elements have different spectra.

Web Links for Part IV of the Report